Chapter 19 “ Acids, Bases, and Salts ”

1. Chapter 19“Acids, Bases, and Salts”

2. Section 19.1Acid-Base Theories • OBJECTIVES: • Define the properties of acids and bases.

3. Section 19.1Acid-Base Theories • OBJECTIVES: • Compare and contrast acids and bases as defined by the theoriesof: a) Arrhenius, b) Brønsted-Lowry, and c) Lewis.

4. Properties of Acids • They taste sour (don’t try this at home). • They can conduct electricity. • Can be strong or weak electrolytes in aqueous solution • React with metals to form H2 gas. • Change the color of indicators (for example: blue litmus turns to red). • React with bases (metallic hydroxides) to form water and a salt.

5. Properties of Acids • They have a pH of less than 7 (more on this concept of pH in a later lesson) • They react with carbonates and bicarbonates to produce a salt, water, and carbon dioxide gas • How do you know if a chemical is an acid? • It usually starts with Hydrogen. • HCl, H2SO4, HNO3, etc. (but not water!)

6. Acids Affect Indicators, by changing their color Blue litmus paper turns red in contact with an acid (and red paper stays red).

7. Acids have a pH lessthan 7

8. Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas: HCl(aq) + Mg(s)→ MgCl2(aq) + H2(g) This is a single-replacement reaction

9. Acids React with Carbonates and Bicarbonates HCl + NaHCO3 Hydrochloric acid + sodium bicarbonate NaCl + H2O + CO2 salt + water + carbon dioxide An old-time home remedy for relieving an upset stomach

10. Effects of Acid Rain on Marble(marble is calcium carbonate) George Washington: BEFORE acid rain George Washington: AFTER acid rain

11. Acids Neutralize Bases HCl + NaOH → NaCl + H2O -Neutralization reactions ALWAYS produce a salt (which is an ionic compound) and water. -Of course, it takes the right proportion of acid and base to produce a neutral salt

12. Sulfuric Acid = H2SO4 • Highest volume production of any chemical in the U.S. (approximately 60 billion pounds/year) • Used in the production of paper • Used in production of fertilizers • Used in petroleum refining; auto batteries

13. Nitric Acid = HNO3 • Used in the production of fertilizers • Used in the production of explosives • Nitric acid is a volatile acid – its reactive components evaporate easily • Stains proteins yellow (including skin!)

14. Hydrochloric Acid = HCl • Used in the “pickling” of steel • Used to purify magnesium from sea water • Part of gastric juice, it aids in the digestion of proteins • Sold commercially as Muriatic acid

15. Phosphoric Acid = H3PO4 • A flavoring agent in sodas (adds “tart”) • Used in the manufacture of detergents • Used in the manufacture of fertilizers • Not a common laboratory reagent

16. Acetic Acid = HC2H3O2 (also called Ethanoic Acid, CH3COOH) • Used in the manufacture of plastics • Used in making pharmaceuticals • Acetic acid is the acid that is present in household vinegar

17. Properties of Bases(metallic hydroxides) • React with acids to form water and a salt. • Taste bitter. • Feel slippery (don’t try this either). • Can be strong or weak electrolytes in aqueous solution • Change the color of indicators (red litmus turns blue).

18. Examples of Bases(metallic hydroxides) • Sodium hydroxide, NaOH(lye for drain cleaner; soap) • Potassium hydroxide, KOH (alkaline batteries) • Magnesium hydroxide, Mg(OH)2(Milk of Magnesia) • Calcium hydroxide, Ca(OH)2(lime; masonry)

19. Bases Affect Indicators Red litmus paper turns blue in contact with a base (and blue paper stays blue). Phenolphthalein turns purple in a base.

20. Bases have a pH greaterthan 7

21. Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH)2 Magnesium salts can cause diarrhea (thus they are used as a laxative) and may also cause kidney stones. MgCl2 + 2 H2O

22. Naming Binary Acids (2 elements) Start with “hydro”, then end with the anion (-ic ending) Ex) HCl = hydrochloric acid How about HBr? What’s the formula for hydrosulfuric acid?

23. Polyatomic acids Contain H’s AND a polyatomic anion Names do NOT start with “hydro” Named after their anion (if it has a sulfate the acid name becomes sulfuric acid.) Remember “If you ATE an acid that would be ICKY”!! Charges must balance. Ex) H2SO4 = Sulfuric Acid H3PO4 =

24. Polyatomic Acids continued If the anion is ends with -ite (sulfite, nitrite) use the same rules as –ate acids, but end them with –ous acid. Ex) HNO2 = Nitrous Acid H2SO3 =

25. Naming Bases Just follow the regular rules for naming ionic compounds: Metal (cation) first, followed by nonmetal (anion)…hydroxide! Balance charges NaOH = Sodium hydroxide Mg(OH)2 = LiOH=

26. Section 19.2Hydrogen Ions and Acidity • OBJECTIVES: • Describe how [H1+] and [OH1-] are related in an aqueous solution.

27. Section 19.2Hydrogen Ions and Acidity • OBJECTIVES: • Classify a solution as neutral, acidic, or basic given the hydrogen-ion or hydroxide-ion concentration.

28. Section 19.2Hydrogen Ions and Acidity • OBJECTIVES: • Convert hydrogen-ion concentrations into pH values and hydroxide-ion concentrations into pOH values.

29. Section 19.2Hydrogen Ions and Acidity • OBJECTIVES: • Describe the purpose of an acid-base indicator.

30. Hydrogen Ions from Water • Water ionizes, or falls apart into ions: H2O ↔ H1+ + OH1- • Called the “self ionization” of water • Occurs to a very small extent: [H1+ ] = [OH1-] = 1 x 10-7 M • Since they are equal, a neutral solution results from water Kw = [H1+ ] x [OH1-] = 1 x 10-14M2 • Kw is called the “ion product constant” for water

31. Ion Product Constant • H2O ↔ H1+ + OH1- • Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x 10-14 M2 • If [H+] > 10-7 then [OH-] < 10-7 • If [H+] < 10-7 then [OH-] > 10-7 • If we know one, other can be determined • If [H+] > 10-7, it is acidic and [OH-] < 10-7 • If [H+] < 10-7, it is basic and [OH-] > 10-7 • Basic solutions also called “alkaline”

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33. The pH concept – from 0 to 14 • pH = pouvoir hydrogene (Fr.) “hydrogen power” • definition: pH = -log[H+] • in neutral pH = -log(1 x 10-7) = 7 • in acidic solution [H+] > 10-7 • pH < -log(10-7) • pH < 7 (from 0 to 7 is the acid range) • in base, pH > 7 (7 to 14 is base range)

35. Calculating pOH • pOH = -log [OH-] • [H+] x [OH-] = 1 x 10-14 M2 • pH + pOH = 14 • Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid • Not greatly used like pH is.

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38. Measuring pH • Why measure pH? • Everyday solutions we use - everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc. • Sometimes we can use indicators, other times we might need a pH meter

39. How to measure pH with wide-range paper 1. Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod. 2.Compare the color to the chart on the vial – then read the pH value.

40. Some of the many pH Indicators and theirpH range

41. Acid-Base Indicators • Although useful, there are limitationsto indicators: • usually given for a certain temperature (25 oC), thus may change at different temperatures • what if the solution already has a color, like paint? • the ability of the human eye to distinguish colors is limited